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CYANAMIDE

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Originally appearing in Volume V19, Page 716 of the 1911 Encyclopedia Britannica.
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CYANAMIDE . The See also:

combination of See also:nitrogen with See also:oxygen was first effected by See also:Cavendish in 1785, who employed a spark See also:discharge. The See also:process was See also:developed by Madame Lefebre in 1859; by Meissner in 1863, who found that moist gases gave a better result; and by See also:Prim in 1882, who sparked the gases under pressure; it was also used by See also:Lord See also:Rayleigh in his See also:isolation of See also:argon (q.v.).. It was not, however, a commercial success, and the same result attended See also:Siemens and Halske's application of the silent discharge. More effective was the electric arc. In 1892 See also:Sir W. See also:Crookes showed that the arc brought about combination; and in 1897 Lord Rayleigh went into the process more fully. But the first careful working-out of the conditions was made in 1900 by A. McDougall and F. Howles, who, employing a high tension alternating arc, showed that the effectiveness depended upon the temperature. The commercial manufacture of nitric See also:acid was attempted by C. S.

See also:

Bradley and D. R. Lovejoy at See also:Niagara Falls, who passed atmospheric See also:air, or air enriched with oxygen, about a high tension arc made as See also:long as possible; but the See also:company (the Atmospheric Products Company) was a failure. Better results have attended the process of K. Birkeland and S. Eyde, which is being worked on a large See also:scale at Notodden, See also:Norway. The arc is produced by leading a current of about 5000 volts equatorially between the poles of an electromagnet; this produces what is practically a disk of See also:flame, 62 ft. in See also:diameter and having a temperature of about 3000°. The disk really consists of a See also:series of successive arcs which increase in See also:size until they burst. The first product of the reaction is nitric See also:oxide, which on cooling with the residual gases produces nitrogen peroxide. The cooled gases are then led into towers where they meet a stream of See also:water coming in the contrary direction. Nitric acid (up to 5o%) is formed in the first See also:tower, and weaker acids in the successive ones; the last tower contains See also:milk of See also:lime which combines with the gases to See also:form See also:calcium nitrite and nitrate (this product, being unsuitable as a manure, is decomposed with the acid, and the evolved gases sent back). It was found advantageous not to See also:work for acid but for a basic calcium nitrate (normal calcium nitrate being very deliquescent); for this purpose the acid is treated with the requisite amount of milk of lime.

In the process of the Badische Anilin- and Soda-Fabrik, the arc, which is said to be 30 to 50 ft. long or more, is formed in a long See also:

tube, and the gases are sent See also:round the arc by obliquely injecting them. A 30% acid is said to be formed. I. Moscicki and J. von Kowalski have patented a process wherein the arc is formed at two See also:vertical concentric See also:copper electrodes and rotated by an electromagnet; it is worked at See also:Vevey, See also:Switzerland. The Rankin process, of which very little is known, produces the arc with much See also:lower current. The See also:conversion of nitrogen into See also:ammonia by See also:electricity has received much See also:attention, but the commercial aspect appears to have been first worked out by de Hemptinne in 1900, who used both the spark and silent discharge on mixtures of See also:hydrogen and nitrogen, and found that the pressure and temperature must be kept See also:low and the spark See also:gap narrow. J. Schlutius in 1903 employed Dowson See also:gas as a source of hydrogen, and induced combination by means of See also:platinum and the silent discharge. Several non-See also:electrical processes have been devised. In 1862 Fleck passed a mixture of See also:steam, nitrogen and See also:carbon monoxide over red-hot lime, whilst in 1904 Woltereck induced combination by passing steam and air over red-hot See also:iron oxide (See also:peat is used in practice). In de Lambilly's process air and steam is led over See also:white-hot See also:coke, and carbon dioxide or monoxide removed from the escaping gases according as ammonium formate or carbonate is wanted. The residual gas is then passed through a tube containing porous materials, such as See also:wood- or See also:bone-See also:charcoal, platinized See also:pumice or spongy platinum, then mixed with steam and again forced through the tube.

The reactions are represented as (I) N2+3H2+2CO ±2H20=2H•CO2NH4 (Ammonium formate). (2) N2+3H2+2CO2-f2H20=2HO•CO2NH4 (Ammonium carbonate). The best temperature for the first reaction is between 8o°C. and 13o°C. and for the second between 4o C. and 6o° C. In another process, which originated with C. Kaiser (Abst. J.C.S., 1907, ii. p. 862), calcium is heated in a current of hydrogen, and nitrogen passed over the hydride so formed; this gives ammonia and calcium nitride, the latter of which gives up its nitrogen as ammonia and reforms the hydride when heated in a current of hydrogen. The fixation of nitrogen as a nitride has not been attended with commercial success. H. Mehner patented See also:

heating the oxides of See also:silicon, See also:boron or See also:magnesium with See also:coal or coke in an electric See also:furnace, and then passing in nitrogen, which forms, with the See also:metal liberated by the See also:action of the carbon, a readily decomposable nitride. For an extended bibliography see Bulletin No. 63 of the See also:Bureau of Soils, U.S.

See also:

Department of See also:Agriculture (See also:Washington, 1910). Nitrogen is a colourless, tasteless and odourless gas, which is only very slightly soluble in water. It is slightly lighter than air. Lord Rayleigh in 1894 found that the See also:density of atmospheric nitrogen was about z % higher than that of chemically prepared nitrogen, a See also:discovery which led to the isolation of the rare gases of the See also:atmosphere (see ARGON). The values obtained are shown below. Atmospheric Chemical Nitrogen. Nitrogen. 0.97209 0.96727 Lord Rayleigh, Chem. See also:News, 1897, 76, P. 315. 0.9720 0.9671 A. Leduc, Comptes rendus, 1896, 123, p.

805. Nitrogen is a very inert gas: it will neither See also:

burn nor support the See also:combustion of See also:ordinary combustibles. It combines directly with See also:lithium, calcium and magnesium when heated, whilst nitrides of the rare See also:earth metals are also produced when their oxides are mixed with magnesium and heated in a current of nitrogen (C. Matignon, Comptes rendus, 1900, 131, p. 837). Nitrogen has been liquefied, the See also:critical temperature being -149° C. and the critical pressure 27.54 atmospheres. The liquefied gas boils at -195.5° C., and its specific gravity at its boiling point is 0.8103 (E. C. C. Baly and F. G. Donnan, Jour.

Chem. See also:

Soc., 1902, 81, p. 912). Compounds. Nitrogen combines with hydrogen to form ammonia, NH3, See also:hydrazine, N2H4, and See also:azoimide, N3H (qq.v.); the other known hydrides, N4H4 and N5H5, are salts of azoimide, viz. NH4.N3 and N2H4•N3H. Nitrogen trichloride, NCl3, discovered by P. L. See also:Dulong in 1811 (Schweigg. Journ., 1811, 8, p. 302), and obtained by the action of See also:chlorine or See also:sodium hypochlorite on ammonium chloride, or by the See also:electrolysis of ammonium chloride See also:solution, is a very volatile yellow oil. It possesses an extremely pungent See also:smell, and its vapour is extremely irritating to the eyes.

It is a most dangerous explosive(see D. L. See also:

Chapman and L. Vodden, Jour. Chem. Soc., 1909, 95, p. 138). Chlorine azide, C1•N3, was discovered by F. Raschig in 1908 (see Az0IMIDE) ; the corresponding See also:iodine See also:compound had been obtained in 1900 by A. Hantzsch (Ber., 33, p. 522). For the so-called nitrogen iodide see AMMONIA.

Nitrogen forms five oxides, viz. nitrous oxide, N20, nitric oxide, NO, nitrogen trioxide, N203, nitrogen peroxide, NO2, and nitrogen pentoxide, N2O5, whilst three oxyacids of nitrogen are known: yponitrous acid, H2N202, nitrous acid, HNO2, and nitric acid. HNO3 (q.v.). The first four oxides are gases, the fifth is a solid. Nitrous oxide, N20, isolated in 1776 by J. See also:

Priestley, Who obtained it by reducing nitrogen peroxide with iron, may be prepared by heating ammonium nitrate at 17o-26o° C., or by reducing a mixture of nitric and sulphuric acid with See also:zinc. It is a colourless gas, which is practically odourless, but possesses a sweetish See also:taste. It is some-what soluble in water. When liquefied it boils at -89.8° C., and by further cooling may be solidified, the solid melting at -1o2.3° C (W. See also:Ramsay, Chem. News, 1893, 67, p. 140). It does not burn, but supports the combustion of heated substances almost as well as qxygen.

It is used as an anaesthetic, principally in See also:

dentistry, producing when inhaled a See also:condition of hysterical excitement often accompanied by loud See also:laughter, whence it is sometimes called " laughing gas." Nitric oxide, NO, first obtained by See also:Van See also:Helmont, is usually pre-pared by the action of dilute nitric acid (sp. gr. 1.2) on copper. This method does not give a pure gas, varying amounts of nitrous oxide and nitrogen being See also:present (see NIT Ric Acxu). In a purer condition it may be obtained by the action of sulphuric acid on a mixture of See also:potassium nitrate and ferrous sulphate, or of hydrochloric acid on a mixture of potassium nitrate and ferric chloride. It is also formed by the action of concentrated sulphuric acid on sodium nitrite in the presence of See also:mercury. It is a colourless gas which is only sparingly soluble in water. It may be liquefied, its critical temperature being -93.5°, and the liquid boils at -153.6° C. It is not a supporter of combustion, unless the sustance introduced is at a sufficiently high temperature to decompose the gas, when combustion will continue at the expense of the liberated oxygen. If the gas be mixed with the vapour of carbon disulphide, the mixture See also:burns with a vivid See also:lavender-coloured flame. Nitric oxide is soluble in solutions of ferrous salts, a dark See also:brown solution being formed, which is readily decomposed by See also:heat, with See also:evolution of nitric oxide. It combines with oxygen to form nitrogen peroxide. Nascent hydrogen reduces it to See also:hydroxylamine (q.v.), whilst solutions of hypochlorites oxidize it to nitric acid.

In some instances it reacts as a reducing See also:

agent, e.g. See also:silver oxide is reduced to metallic silver at 170° C., See also:lead dioxide to the monoxide and See also:manganese dioxide to sesquioxide. Nitrogen trioxide, N2O3, was first mentioned by J. R. See also:Glauber in 1648 as a product of the reaction between nitric acid and arsenious oxide. Sir W. Ramsay (Jour. Chem. Soc., 1890, 5, p. 590), by distil-See also:ling arsenious oxide with nitric acid and cooling the distillate, obtained a See also:green liquid which consisted of nitrogen trioxide and peroxide in varying proportions, and concluded that the trioxide could not be obtained pure. He then tried the See also:direct combination of nitric oxide with liquid nitrogen peroxide. A dark See also:blue liquid is produced, and the first portions of gas boiling off from the mixture correspond fairly closely in See also:composition with nitrogen trioxide. H.

B. See also:

Baker (Jour. Chem. Soc., 1907, 91, p. 1862) obtained nitrogen trioxide in the gaseous form by volatilizing the liquid under See also:special conditions. L. Francesconi and N. See also:Sciacca (Ganz., 1904, 34 (i.), p. 447) have shown that liquid nitric oxide and oxygen, or gaseous nitric oxide and liquid oxygen, mixed in all proportions and yielded nitrogen trioxide, whilst gaseous nitric oxide mixed with excess of oxygen always gave the trioxide if the mixture was kept below —I1o° C. They also See also:state that nitrogen trioxide is See also:stable at ordinary pressure up to -21° C. N. M. v.

Wittorf (Zeit. anorg. Chem., 1904, 41, p. 85) obtained blue crystals of the trioxide (melting at -103° C.) on saturating liquid nitrogen peroxide with nitric oxide and cooling the mixture. The liquid prepared by Baker is green in See also:

colour, and has a specific gravity 1.11 at ordinary temperature, but below -2° C. becomes of a deep See also:indigo blue colour. It forms a See also:mass of deep blue crystals at the temperature of liquid air. It is exceedingly soluble in concentrated sulphuric acid. Nitrogen peroxide, NO2 or N2O4, may be obtained by mixing oxygen with nitric oxide and passing the red gas so obtained through a freezing mixture. The See also:production of this red gas when air is mixed with nitric oxide was mentioned by R. See also:Boyle in 1671. Nitrogen peroxide is also prepared by heating lead nitrate and passing the products of decomposition through a tube surrounded by a freezing mixture, when the gas liquefies. At low temperatures it is a colour-less crystalline solid which melts at -1014° C. (W.

Ramsay, Chem. News, 1900, 61, p. 91). As the temperature increases the liquid becomes yellowish, the colour deepening with rise of temperature until at +15° C. it has a deep See also:

orange tint. The liquid boils at about 22° C. This See also:change of colour is accompanied by a change in the vapour density, and is explained by the fact that nitrogen peroxide consists of a mixture of a colourless compound N2O4, and a red-brown gas NO2, the latter increasing in amount at the expense of the former as the temperature is raised (G. Salet, Comptes rendus, 1868, 67, p. 488; see also E. and L. Natanson, Wied. See also:Ann., 1885, 24, p. 454; 1886, 27, p. 6o6).

M. See also:

Berthelot and J. Ogier (Pull. Soc. Chim., 1882 [21, 37, p. 434; 38, p. 60) have also shown that the specific heat of the gas decreases with increase of temperature until it reaches a minimum at about 198-253° C. Cryoscopic determinations of the molecular See also:weight of nitrogen peroxide dissolved in tlacial acetic acid show that it corresponds to the molecular See also:formula 204 at low temperatures (W. Ramsay, Jour. Chem. Soc., 1888, 53, p. 621).

Nitrogen peroxide is the most stable oxide of nitrogen. It is decomposed by water, giving at o° C. a mixture of nitric and nitrous acids: 2NO2+See also:

H2O=HNO3+HNO2. It combines with sulphuric acid to form nitro-sulphonic acid, S02(OH) (NO2). It does not support the combustion of a See also:taper, but burning See also:phosphorus and red-hot carbon will continue to burn in the gas. It converts many metallic oxides into mixtures of nitrates and nitrites, and attacks many metals, forming nitrates and being itself reduced to nitric oxide. It is an energetic oxidizing agent. Nitrogen pentoxide, N205, was first obtained in 1849 by H. Sainte-Claire-Deville (Ann. Chim. Phys., 1850 [31, 28, p. 241) by the action of dry chlorine on silver nitrate: 4AgNO3+2C12=4AgCI+2N205 +02. It may also be obtained by distilling nitric acid over phosphorus pentoxide.

It crystallizes in large prisms which melt at 29-30° C. to a yellowish liquid, which boils at 45-50° C. with rapid decomposition. It is very unstable, decomposing slowly even at ordinary temperatures. It dissolves in water, forming nitric acid. Hyponitrous acid, H2N202, was first obtained in the form of its salts by E. See also:

Divers in 1871 (Chem. News, 23, p. 206) by reducing a solution of potassium nitrite with sodium See also:amalgam, and subsequent precipitation as silver See also:salt. Hyponitrites also result when hydroxyamido-sulphonates, e.g. HO.NH•S03Na, are hydrolysed by See also:caustic alkalis (E. Divers and T. Haga, Jour. Chem.

Soc., 1889, 55, p. 76o), or when benzsulphohydroxamic acid, C6HES02•NH.0H, is treated in the same manner (0. See also:

Piloty, Ber., 1896, 29, p. 1560). They may also be prepared by the action of mercuric or cupric oxides on alkaline solutions of hydroxylamine (A. Hantzsch, Ann., 1896, 292, p. 317) ; by the action of hydroxylamine sulphate on alkaline nitrites in the presence of lime or calcium carbonate, the mixture being rapidly heated to 6o° C.; or by the See also:hydrolysis of dimethyl nitroso-oxyurea, (See also:CH3)2N•CO•N(NO)•OH (A. Hantzsch, Ber., 1897, 30, p. 2356). The See also:free acid, which crystallizes in brilliant scales, is best prepared by decomposing the silver salt with an ethereal solution of hydrochloric acid. It is very explosive, dissolves readily in water and behaves as a dibasic acid. It does not liberate iodine from potassium iodide, neither does it decolorize iodine solution.

See also:

Bromine oxidizes it to nitric acid, but the reaction is not quantitative. In acid solution, potassium permanganate oxidizes it to nitric acid, but in alkaline solution only to nitrous acid. It decomposes slowly on See also:standing, yielding water and nitrous oxide. The silver salt is a See also:bright yellow solid, soluble in dilute sulphuric and nitric acids, and may be crystallized from concentrated solutions of ammonia. It slowly decomposes on exposure or on heating. The calcium salt, CaN202.4H20, formed by the action of calcium chloride on the silver salt in the presence of a small quantity of nitric acid, is a lustrous crystalline See also:powder, almost insoluble in water but readily soluble in dilute acids. It is decomposed by sulphuric acid, with evolution of nitrous oxide. Nitrous acid, HNO2, is found to some extent in the form of its salts in the atmosphere and in See also:rain water. The pure acid has not yet been obtained, since in the presence of water it decomposes with formation of nitric acid and liberation of nitric oxide: 3HNO2 =HNO3+2NO+H20. Its salts may be obtained in some cases by heating the corresponding nitrates, but the method does not give See also:good results. Sodium nitrite, the most commonly used salt of the acid, is generally obtained by heating the nitrate with metallic lead; by heating sodium nitrate with See also:sulphur and sodium hydroxide, the product then being fractionally crystallized (Read, Holliday & Sons: 3NaNO3+S+2NaOH = Na2SO4+3NaNO2+H20; by oxidizing atmospheric nitrogen in an electric arc, keeping the gases above 300° C., until absorption in alkaline hydroxide solution is effected (See also:German Pat. 188188) ; or by passing air, or a mixture of oxygen and ammonia, over heated metallic oxides (ibid., 168272).

The salts of the acid are colourless or faintly yellow. In aqueous solution the free acid acts as an oxidizing agent, See also:

bleaching indigo and liberating iodine from potassium iodide, or it may See also:act as a reducing agent since it readily tends to pass into nitric acid: consequently it discharges the colour of acid solutions of permanganates and chromates. The acid finds considerable use in organic See also:chemistry, being employed to discriminate between the different types of See also:alcohols and of See also:amines, and also in the production of diazo, See also:azo and diazo-amino compounds. It may be recognized by the blue colour it gives with diphenylamine sulphate and by its reaction with potassium iodide-See also:starch See also:paper. Nitrosyl chloride, NOCI, is obtained by the direct See also:union of nitric oxide with chlorine; or by distilling a mixture of concentrated nitric and hydrochloric acids, passing the resulting gases into concentrated sulphuric acid and heating the so-formed nitrosyl hydrogen sulphate with dry salt: HNO3+3HC1=NOC1+C12 +H20; NOCI + H2SO4 = HCI + NO•SO4H; NO. SO4H + NaCl =NOCI+NaHSO4 (W. A. See also:Tilden, Jour. Chem. Soc., 186o, p. 63o). It is also prepared by the action of phosphorus pentachloride on potassium nitrite or on nitrogen peroxide.

It is an orange-coloured gas which may be readily liquefied and by further cooling may besolidified. The liquid boils at -5° C. and the solid melts at -65° C. It forms See also:

double compounds with many metallic chlorides, and finds considerable application as a means of separating various members of the terpene See also:group of compounds. It is readily decomposed by water and alkaline hydroxides, yielding a mixture of nitrite and chloride. On treatment with silver fluoride it yields nitrosyl fluoride, NOF (O. See also:Ruff, Zeit. anorg. Chem., 1905, 47, p. 190). Nitroxyl fluoride, NO2F, is formed by the action of See also:fluorine on nitric oxide at the temperature of liquid oxygen (H. See also:Moissan and P. See also:Lebeau, Comptes rendus, 1905, 140, pp. 1573, 1621).

It is a gas at ordinary temperature; when liquefied it boils at -63.5° C. and on solidification melts at -139° C. Water decomposes it into nitric and hydrofluoric acids. Nitramide, NH2NO2, is obtained by the action of sulphuric and nitric acids on potassium imidosulphonate, or by the action of See also:

ice-See also:cold sulphuric acid on potassium nitro-carbamate (J. Thiele and A. See also:Lachmann, Ann., 1895, 288, p. 297): NO2•NK-CO2K+H2SO4 =NH2NO2+K2SO4+See also:CO2. It crystallizes in prisms or leaflets which melt at 72-75°C. and are readily soluble in water and in all organic solvents except ligroin. It is somewhat volatile at ordinary temperature, and its aqueous solution possesses a strongly acid reaction. It is very unstable, decomposing.into nitrous oxide and water when mixed with copper oxide, lead chromate or even powdered See also:glass. On reduction it gives a strongly reducing substance, probably hydrazine. According to A. Hantzsch (Ann., 1896, 292, pp.

340 et seq.) hyponitrous acid and nitramide are to be regarded as stereo-isomers, being the See also:

anti- and syn- forms of the same compound. Thiele, however, regards nitramide as imidonitric acid, HN :NO(OH). Nitrogen sulphide, N4S4, first obtained by W. See also:Gregory (Jour. pharm., 1835, 21, p. 315) by the action of ammonia on sulphur chloride, has been investigated by O. Ruff and E. Geisel (Ber., 1904, 37, p. 1573; 2905, 38, p. 2659), who also obtained it by dissolving sulphur in liquid ammonia. It is a reddish-yellow crystalline solid, insoluble in water and melting at 178° C. It explodes readily when melted or subjected to See also:shock. Dry hydrochloric acid gives ammonia but no nitrogen; with ammonia it gives N:SNH2 and S :S(See also:NH2)2 ; and with secondary amines it forms thiodiamines, S(NR2)2, nitrogen and ammonia being liberated.

When heated with CS2 to See also:

loo' C. under pressure, it forms liquid nitrogen sulphide, N2S,,a See also:mobile red liquid which solidifies to an iodine-like mass of crystals which melt at lo-11° C. Water, alkalis and acids decompose it into sulphur and ammonia (W. Muthmann, Zeit. anorg. Chem., 1897, 13, p. 200). For sulphonic acids containing nitrogen see AMMONIA. Numerous determinations of the atomic weight of nitrogen have been made by different observers, the values obtained varying some-what according to the methods used. These methods have been purely chemical (either See also:gravimetric or volumetric), See also:physical (determinations of the density of nitrogen, nitric oxide, &c.) or physicochemical. P. A. Guye has given a critical discussion of the relative accuracy of the gravimetric and physico-chemical methods, and favours the latter, giving for the atomic weight a value less than 14.01. The more important papers dealing with the subject are: J.

See also:

Stas, fEuvres completes, i. pp. 342 et seq.; Lord Rayleigh, Proc. See also:Roy. Soc. (1894), 55, p. 340; (1904) 73, p. 153; G. See also:Dean, Jour. Chem. Soc. (190,), 79, p. 147; R.

W. See also:

Gray, Jour. Chem. Soc. (,906), 88, p. 1174; A. See also:Scott, Proc. Chem. Soc. (1905), 21, p. 309; P. A.

Guye, Chem. News (1905), 92, pp. 261 et seq.; (1906) 93, p. 13 et seq. ; D. Berthelot, Comptes rendus (1907), 144, p. 269.

End of Article: CYANAMIDE

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